Introduction to energetics and enthalpy

In this lesson, we will learn:

• How energy changes are measured in chemical reactions.
• The standard conditions at which enthalpy changes are reported.
• To construct enthalpy diagrams for endothermic and exothermic reactions.
• The definitions of enthalpy of reaction, formation, combustion and neutralization.

• Outside of chemistry, to be energetic means to have energy. In chemistry, energetics is about changes in amount of energy in chemical reactions, going from reactants to products.
This is useful because many chemical reactions give out or absorb energy as heat that we can measure.
The net amount of heat energy absorbed or released from a reaction at constant pressure, is called an enthalpy change, symbol $\Delta H$ (delta H).
• Measuring and reporting enthalpy changes is done at standard conditions. This is so we have a reference point where different data sets can be easily compared:
• Standard conditions define pressure at 100kPa and 1M concentration for solutions.
• Temperature is not actually part of standard conditions, but for convenience and fair comparison, most data collected is at 298 K (25°C).
• Standard conditions is not the same as standard temperature and pressure (STP), which does strictly define 100 kPa and 273 K!
• The units of enthalpy change are kJ mol^-1 (kilojoules per mole). This is measuring energy in kJ per one mole of substance produced or reacted.
• The standard enthalpy change of reaction, $\Delta H^{\ominus} _{r}$, is the enthalpy change measured when a reaction takes place at standard conditions with all reactants and products in their standard states. Standard state is the common stable state that you find substances in at standard conditions (such as H₂O as liquid water, not as steam/ice).


• The enthalpy change of a reaction comes from two sources. Heat energy is in one of two places: inside the molecules that are reacting (the system), or outside the system, in the rest of the universe (the surroundings).
  When we put it this way, energy is transferred during two events in the reaction:
• To begin a chemical reaction, heat energy must be supplied by the surroundings to break bonds in the reactant molecules. The stronger the bond, the more energy is needed.
• For example, the N≡N triple bond is a lot stronger than the Cl-Cl bond, so more energy is needed to break the N≡N bond than the Cl-Cl bond.
• When new bonds form to make the products, energy is released back into the surroundings. The more stable/stronger the bond, the more energy is released when it forms.
• For example, the strong N≡N triple bond releases a lot more energy when it forms as a product of a reaction than when the Cl-Cl bond forms. Forming the strong N≡N bond releases a lot of energy into the environment, so the N≡N molecule itself is in a lower energy state.
• This is why chemists call stable molecules low energy they have released most of their energy to the environment. In the same way, unstable molecules are sometimes called high energy.

Total energy is conserved in chemical reactions. This means no heat energy is created or destroyed; it is only transferred into or out of the system.


• Its very rare for a reaction to need a certain amount of energy to break up the reactants, and then give back the exact same amount of energy when products form.
  Almost all chemical reactions have an overall effect of absorbing energy from or releasing energy into the surroundings. There are two definitions for this:
• An exothermic reaction is a chemical reaction that has the overall effect of releasing heat energy to the surroundings.
• This means that the energy released when forming product bonds was greater than the energy required when breaking up the reactant bonds.
• Exothermic enthalpy changes are always given a negative sign, to show that the chemical substances have become lower energy (the energy they released is now in the surroundings).
• An endothermic reaction is a reaction that has the overall effect of absorbing heat energy from the surroundings.
• This means that the energy required to break up the reactant molecules was greater than the energy released when the new product bonds formed.
• Endothermic enthalpy changes are always given a positive sign, to show that the chemical substances have become higher energy (because they absorbed more energy from the surroundings as reactants than they released when they became products).


• Enthalpy is a measure of heat energy, so the enthalpy change of a reaction can be seen in a temperature change of the surroundings:
• An exothermic reaction will increase the temperature of the surroundings. For example, an exothermic dissolving process will cause the temperature of the water solvent to increase.
• An endothermic reaction will increase the temperature of the surroundings. For example, an endothermic dissolving process will reduce the temperature of the water solvent.


• These two types of reactions can be shown by drawing an enthalpy (energy level) diagram.
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• There are some specific enthalpy definitions that should be known for common chemical changes:
• The standard enthalpy of combustion, $\triangle H^{\ominus}$_c, is the enthalpy change measured when one mole of a substance is completely reacted with oxygen at standard conditions.
• The standard enthalpy of neutralization, $\triangle H^{\ominus}$_n, is the enthalpy change measured when an acid and base react to form one mole of water in a neutralization reaction at standard conditions.
• The standard enthalpy of formation, $\triangle H^{\ominus}$_f, is the enthalpy change when one mole of a substance is formed from its constituent elements in their standard states at standard conditions.
These definitions have several things in common, which is exactly the point.
• The definitions are all measuring energy change for one mole of a substance.
• The definitions are all measurements at standard conditions and substances in their standard states.
This is all done so scientists can compare different data sets properly. If the same temperature (298K), same pressure (100kPa) and same amount of substance (1 mole) are all used, then we can prove that it is the different substances that produce the different enthalpy changes.


• Enthalpy changes of a reaction is measured in several ways:
• We can apply Hesss law. In short, because heat energy is never created or destroyed in reactions, the enthalpy change from reactants A to products B is the same regardless of which route or set of steps are taken to get there.
• We can use mean bond enthalpy. In short, because a reaction is just breaking reactant bonds and making product bonds, if we know the amount of energy (in kJ mol-1) needed to do that and the molar ratios in the reaction, we can find out the total amount of heat energy exchanged.
• We can use calorimetry, where the release of heat energy to the surroundings can be measured by the temperature increase.

We will see methods to calculate using all three of these in this chapter.